Chapter 7 Acids and Bases

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1 Chapter 7 Acids and Bases 7.1 The Nature of Acids and Bases 7.2 Acid Strength 7.3 The ph Scale 7.4 Calculating the ph of Strong Acid Solutions 7.5 Calculating the ph of Weak Acid Solutions 7.6 Bases 7.7 Polyprotic Acids 7.8 Acid-Base Properties of Salts 7.9 Acid Solutions in Which Water Contributes to the H + Concentration 7.10 Strong Acid Solutions in Which Water Contributes to the H + Concentration 7.11 Strategy for solving Acid-Base Problems: A Summary Acids and Bases Acids were first recognized as substances that taste sour. sulfuric acid reacts with water in your skin a very exothermic reaction Saponification Bases taste bitter, feel slippery. lye (sodium hydroxide) reacts with lipids and oils in skin to produce a soap molecule ph of foods: 1

2 A & B: Arrhenius Theory Acid = a substance containing H + and whose aqueous solution contains more H + ions than OH - ions. Dissolve HA in water... HO 2 HA( aq) H + ( aq) + A ( aq) Base = a substance containing OH - and whose aqueous solution contains more OH - ions than H + ions. Dissolve BOH in water... + BOH( aq) B ( aq) + OH ( aq) Neutralization: produces water and an aqueous salt. H A( aq) + BOH( aq) AB( aq) + H OH( aq) Net ionic equation: H ( aq) + OH ( aq) H OH( aq) A & B: Bronsted-Lowry Theory Bronsted-Lowry: acid/base reactions are protontransfer processes. acid is proton-donor (H + ion donor). base is proton- acceptor (H + ion acceptor). B + H A B H + A When an acid gives its proton to water, water is acting as a base. H 2 O + H A H 2 O H + A - 2

3 Arrhenius vs. Bronsted-Lowry Arrhenius: acid... [H + ] > [OH - ] base... [H + ] < [OH - ] Bronsted-Lowry Bronsted-Lowry: proton transfer: from an acid (protondonor), to a base (proton acceptor) to form a covalent bond. the base donates the electrons for the bond the proton accepts the electrons for the bond Arrhenius HA + B A + H B + A & B: Lewis Acids BF 3 NH 3 F 3 B-NH 3 Lewis Acid: electron pair acceptor Lewis Base: electron pair donator Lewis acid/base reactions are not limited to aqueous environments and can have other e-pair acceptors than just H +. 3

4 A & S: Lewis Acids (cont) The formation of a covalent bond is the driving force in all three acid/base schemes. Lewis acid/base theory is the most general. Lewis Bronsted-Lowry Arrhenius That said, we are going to focus on the Bronsted-Lowry definition of acids and bases. A & B: Summary Lewis A + B A B Bronsted-Lowry A H( aq) + B( aq) A ( aq) + H B( aq) Arrhenius A H( aq) + OH ( aq) A ( aq) + H OH( l) 4

5 B L: Conjugate Acid/Base Pairs conjugate pair B + H A B H + A Base Acid Acid Base conjugate pair Every acid has a conjugate base. These species differ by H +. Every base has a conjugate acid. These species differ by H +. Acid Strength HA(aq) + H 2 O(l) H 3 O + (aq) + A - (aq) Acids are quantified by the extent of H 3 O + production. What does H 3 O + look like? We tend to use H + and H 3 O + synonymously. 5

6 BL: Conjugate Acid/Base Pairs (cont) In this example, acetic acid donates a proton to water, leading to the formation of hydronium ion and acetate ion. H + donor (acid) H + acceptor (base) New acid formed New base formed BL: Conjugate Acid/Base Pairs (cont) conjugate pair conjugate pair Base Acid Acid Base p

7 Molecular Model of Conjugate Acids & Bases The reaction of an acid HA with water to form H 3 O+ and a conjugate base HA A Acid Base Conjugate Conjugate acid base Conjugate acid is formed when a proton is transferred to the base. Conjugate base is everything that remains of the acid molecule after a proton is lost. Conjugated acid-base pair consists of two substances related to each other by donating and accepting a single proton. Acid Strength Strong Acid HA HA(aq) + H 2 O(l) H + A - H 3 O + (aq) + A - (aq) Forward reaction dominates A - is a much weaker base than H 2 O HA is a much stronger acid than H 3 O + [H + ] = [A - ] = [HA] init Weak Acid HA HA H + A - Reverse reaction dominates A - is a much stronger base than H 2 O HA is a much weaker acid than H 3 O + [H + ] = [A - ] << [HA] 7

8 K a Acid Strength (cont) HA(aq) + H 2 O(l) + H3O A 1 M 1 M HA 1 M H 3 O + (aq) + A - (aq) The magnitude of K a is a measure of how likely the acid is to dissociate in water. The Acid-Dissociation Constant (K a ) Strong acids dissociate completely in water to form ions: HA (aq) + H 2 O (l) H 3 O + (aq) + A - (aq) In a dilute solution of a strong acid, almost no HA molecules exist: [H 3 O + ] = [HA] init or [HA] eq = 0 K = K a = [H 3O + ][A - ] [HA] at equilibrium, and K a >> 1 Nitric acid is an example: HNO 3 (aq) + H 2 O (l) Weak acids dissociate very slightly into ions in water: HA (aq) + H 2 O (l) H 3 O + (aq) + A - (aq) H 3 O + (aq) + NO 3 - (aq) In a dilute solution of a weak acid, the great majority of HA molecules are not dissociated: [H 3 O + ] << [HA] init or [HA] eq = [HA] init K = K a = [H 3 O + ][A - ] [HA] at equilibrium, and K a << 1 8

9 The Meaning of K a, the Acid Dissociation Constant For the ionization of an acid, HA: HA (aq) + H 2 O (l) H 3 O + (aq) + A - (aq) Therefore: K a = [H 3 O + ] [A - ] [HA] The stronger the acid, the higher the [H 3 O + ] at equilibrium, and the larger the K a : Stronger acid higher [H 3 O + ] larger K a For a weak acid with a relative high K a (~10-2 ), a 1 M solution has ~10% of the HA molecules dissociated. For a weak acid with a moderate K a (~10-5 ), a 1 M solution has ~ 0.3% of the HA molecules dissociated. For a weak acid with a relatively low K a (~10-10 ), a 1 M solution has ~ 0.001% of the HA molecules dissociated. Acid Strength (cont) HA(aq) + H 2 O(l) H 3 O + (aq) + A - (aq) The weaker an acid, the stronger its conjugate base (and vice versa). The conjugate base of an acid wants to get a proton, but it must want it more than H 2 O to be considered a strong conjugate base. The conjugate acid of a base wants to lose a proton, but it must want to lose it more than H 3 O + to be considered a strong conjugate acid. Relative acid strength very strong strong weak very weak Relative conj. base strength very weak weak strong very strong 9

10 Acid Strength (cont.) Which of the following bases would have the weakest conjugate acid: OC 6 H 5 C 2 H 3 O 2 OCl NH 3 Selected Acids and Bases Acids Bases Strong: H + (aq) + A - (aq) Strong: M + (aq) + OH - (aq) hydrochloric, HCl lithium hydroxide, LiOH hydrobromic, HBr sodium hydroxide, NaOH hydroiodoic, HI potassium hydroxide, KOH nitric acid, HNO 3 calcium hydroxide, Ca(OH) 2 sulfuric acid, H 2 SO 4 strontium hydroxide, Sr(OH) 2 perchloric acid, HClO 4 barium hydroxide, Ba(OH) 2 *(M is Group I or II metal) Weak Weak hydrofluoric, HF ammonia, NH 3 phosphoric acid, H 3 PO 4 accepts proton from water to make acetic acid, CH 3 COOH NH 4+ (aq) and OH - (aq) (or HC 2 H 3 O 2 ) 10

11 Classifying the Relative Strengths of Acids and Bases I Strong acids. There are two types of strong acids: 1. The hydrohalic acids: HCl, HBr, and HI 2. Oxyacids in which the number of O atoms exceeds the number of ionizable H atoms by two or more, such as HNO 3, H 2 SO 4, HClO 4 Weak acids. There are many more weak acids than strong ones. Four types, with examples, are: 1. The hydrohalic acid, HF 2. Those acids in which H is NOT bound to O or to a halogen, such as HCN and H 2 S 3. Oxyacids in which the number of O atoms equals or exceeds by one the number of ionizable H atoms, such as HClO, HNO 2, and H 3 PO 4 4. Organic acids (general formula RCOOH), such as CH 3 COOH and C 6 H 5 COOH Classifying the Relative Strengths of Acids and Bases - II Strong bases. Soluble compounds containing O 2- or OH - ions are strong bases. The cations are usually those of the most active metals: 1) M 2 O or MOH, where M= Group 1A(1) metals (Li, Na, K, Rb, Cs) 2) MO or M(OH) 2, where M = Group 2A(2) metals (Ca, Sr, Ba) (MgO and Mg(OH) 2 are only slightly soluble, but the soluble portion dissociates completely.) Weak bases. Many compounds with an electron-rich nitrogen are weak bases (none are Arrhenius bases). The common structural feature is an N atom that has a lone electron pair in its Lewis structure. 1) Ammonia (:NH 3 ) 2) Amines (general formula RNH 2, R 2 NH, R 3 N), such as CH 3 CH 2 NH 2, (CH 3 ) 2 NH, (C 3 H 7 ) 3 N, and C 5 H 5 N : : : : : : : 11

12 12

13 Expressing Acidity: ph [H 3 O + ] in pure water = 1 x 10-7 M [H 3 O + ] in acid rain = 1 x 10-4 M Define power of Hydrogen (ph) ph = -log 10 [H 3 O + ] [H 3 O + ] = 10 -ph Say... Why don t we just use this #? ph represents the order of magnitude of the hydronium ion concentration. A solution at 25 o C with ph = 7.0 means it has the same [H 3 O + ] as pure water, not that there is no acid present. ph = Power of Hydrogen Uses a logarithmic scale (powers of 10) ph units are related to the exponents of [H 3 O + ] ph = -log [H 3 O + ] [H 3 O + ] = 10 -ph [H 3 O + ] in pure water = 1 x 10-7 M ph = -log (1 x 10-7 ) = 7.0 Sig Fig Rule: The number of decimal places in the log is equal to the number of sig figs in the original number. 13

14 ph (cont.) Strong acids can be considered to ionize completely in aqueous solution, increasing [H 3 O + ]. + HA( aq) + H2O( l) A ( aq) + H3O ( aq) We can find ph directly. [H O ] [HA] Strong bases also are considered to completely ionize in aqueous solution resulting in an increase in [OH - ]. + XOH( aq) OH ( aq) + X ( aq) [OH ] [XOH] 0 ph and poh We can define a similar expression for power of hydroxide : poh = -log 10 [OH - ] In aqueous solution at 25 o C ph + poh =

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